1. Aim: How do we perfect our knowledge of the periodic table?
Do now: How is the periodic table arranged?

2. Announcements Flame Test Lab - Thu, Sept. 8 Exam #1- September 12th
Free Response # 2 - due Thu, Sep 15th
Reading Guide #1 and #2 - due September 12th

3. Organizing the Elements
A few elements, such as gold and copper, have been known for thousands of years - since ancient times
Yet, only about 13 had been identified by the year 1700.
As more were discovered, chemists realized they needed a way to organize the elements.

4. Organizing the Elements
Chemists used the properties of elements to sort them into groups.
In 1829 J. W. Dobereiner arranged elements into triads – groups of three elements with similar properties
One element in each triad had properties intermediate of the other two elements

5. Mendeleev’s Periodic Table
By the mid-1800s, about 70 elements were known to exist
Dmitri Mendeleev – a Russian chemist and teacher
Arranged elements in order of increasing atomic mass
Thus, the first “Periodic Table”

6. He left blanks for yet undiscovered elements
Mendeleev
He left blanks for yet undiscovered elements
When they were discovered, he had made good predictions
But, there were problems:
Such as Co and Ni; Ar and K; Te and I

7. A better arrangement
In 1913, Henry Moseley – British physicist, arranged elements according to increasing atomic number
The arrangement used today

9. Another possibility: Spiral Periodic Table

10. Trends in Atomic Size
First problem: Where do you start measuring from?
The electron cloud doesn’t have a definite edge.
They get around this by measuring more than 1 atom at a time.

11. Atomic Size }
Radius
Measure the Atomic Radius - this is half the distance between the two nuclei of a diatomic molecule.

12. ALL Periodic Table Trends
Influenced by three factors:
1. Energy Level
Higher energy levels are further away from the nucleus.
2. Charge on nucleus (# protons)
More charge pulls electrons in closer. (+ and – attract each other)
3. Shielding effect
(blocking effect?)

13. #1. Atomic Size - Group trendsH
As we increase the atomic number (or go down a group). . .
each atom has another energy level,
so the atoms get bigger. Li Na K Rb

14. #1. Atomic Size - Period Trends
Going from left to right across a period, the size gets smaller.
Electrons are in the same energy level.
But, there is more nuclear charge.
Outermost electrons are pulled closer. Na Mg Al Si P S Cl Ar

15. Rb K
Period 2 Na Li
Atomic Radius (pm) Kr Ar Ne H 3 10
Atomic Number

16. #2. Trends in Ionization Energy
Ionization energy is the amount of energy required to completely remove an electron (from a gaseous atom).
Removing one electron makes a 1+ ion.
The energy required to remove only the first electron is called the first ionization energy.

17. Ionization Energy
The second ionization energy is the energy required to remove the second electron.
Always greater than first IE.
The third IE is the energy required to remove a third electron.
Greater than 1st or 2nd IE.

18. Symbol First Second Third
HHeLiBeBCNO F Ne

19. What factors determine IE
The greater the nuclear charge, the greater IE.
Greater distance from nucleus decreases IE
Filled and half-filled orbitals have lower energy, so achieving them is easier, lower IE.
Shielding effect

20. Shielding
The electron on the outermost energy level has to look through all the other energy levels to see the nucleus.
Second electron has same shielding, if it is in the same period

21. Ionization Energy - Group trends
As you go down a group, the first IE decreases because...
The electron is further away from the attraction of the nucleus, and
There is more shielding.

22. Ionization Energy - Period trends
All the atoms in the same period have the same energy level.
Same shielding.
But, increasing nuclear charge
So IE generally increases from left to right.
Exceptions at full and 1/2 full orbitals.

23. Trends in Ionic Size: Cations
Cations form by losing electrons.
Cations are smaller than the atom they came from – not only do they lose electrons, they lose an entire energy level.
Metals form cations.
Cations of representative elements have the noble gas configuration before them.

24. Ionic size: Anions Anions form by gaining electrons.
Anions are bigger than the atom they came from – have the same energy level, but a greater area the nuclear charge needs to cover
Nonmetals form anions.
Anions of representative elements have the noble gas configuration after them.

25. Trends in Electronegativity
Electronegativity is the tendency for an atom to attract electrons to itself when it is chemically combined with another element.
They share the electron, but how equally do they share it?
An element with a big electronegativity means it pulls the electron towards itself strongly!

26. Electronegativity Group Trend
The further down a group, the farther the electron is away from the nucleus, plus the more electrons an atom has.
Thus, more willing to share.
Low electronegativity.

27. Electronegativity Period Trend
Metals are at the left of the table.
They let their electrons go easily
Thus, low electronegativity
At the right end are the nonmetals.
They want more electrons.
Try to take them away from others
High electronegativity.

28. Electron Affinity
The amount of energy needed to remove an electron from a negatively charged ion.
It is also energy released when a single electron is combined with an isolated atom.